The document summarizes several important concepts and laws in chemistry, including: 1) The law of conservation of mass states that matter is neither created nor destroyed in chemical reactions. 2) Laws of definite and multiple proportions describe the ratios in which elements combine in chemical compounds. 3) Gay-Lussac's law of gaseous volumes states that gases combine or are produced in simple volume ratios. 4) Avogadro's law and the mole concept allow chemists to convert between mass and number of particles in chemical reactions.

2. Laws of chemical combination

3. Law of conservation of mass:-
“Law of conservation of mass
states that matter can neither
be created nor be destroyed in
a chemical reaction.”

4. Law of definite proportions
 “In given compound
always contains
exactly the same
proportion of
elements by
weight.”

5. Law of multiple proportions
 “if two elements can combine to
form more than one compound,
the masses of one element that
combine with a fixed mass of
the other element, are in the
ratio of small whole numbers.”

6. Gay lussac’s law of gaseous
volumes
 “when gases combine or are
produced in a chemical
reaction they do so in a
simple ratio by volume
provided all gases are at
same temperature and
pressure.”

7. Avogadro's law
 “Equal volumes of gases contain
equal number of molecules at
standard temperature and
pressure.”

8. Daltons atomic theory
 According to Dalton's atomic theory,
1)Matter consists of indivisible atoms.
2)All the atoms of a given element have identical
properties including identical mass. Atoms of different
elements differ in mass.
3)Compounds are formed when atoms of different
elements combine in a fixed ratio.
4)Chemical reactions involve reorganisation of atoms.
These are neither created nor destroyed in a chemical
reaction.

9. Atomic and molecular
masses
 One atomic mass unit is
defined as a mass exactly
equal to one- twelfth the mass
of one carbon - 12 atom and
1 amu = 1.66056×10–24 g .
 Today, ‘amu’ has been
replaced by ‘u’ which is
known as unified mass.

10. Molecular mass and formula
mass
 Molecular mass is the sum of
atomic masses of the
elements present in a
molecule. It is obtained by
multiplying the atomic mass
of each element by the
number of its atoms and
adding them together.
 Formula unit mass is the
empirical formula of the

11. Mole concept
 One mole is the amount of a
substance that contains as
many particles or entities as
there are atoms in exactly 12 g
of the 12C isotope.
 1 mol of atom = 6.022×1023
entities.
 The mass of one mole of a
substance in grams is called its
molar mass.

12. Limiting reagent
 When any reaction is carried out if
the reactants are not present in the
amounts as required by a balanced
chemical reaction.
 In such situations, one reactant is in
excess over the other. The reactant
which is present in the lesser
amount gets consumed after
sometime and after that no further
reaction takes place whatever be the
amount of the other reactant
present.
 Hence, the reactant which gets
consumed, limits the amount of

13. Concentrations in solutions
 The concentration of a
solution is expressed in 4
different ways.
1) Mass percent
2) Mole fraction
3) Molarity
4) molality

14. Mass percentage
 Mass % of a compound is given
by the formula
 Mass percentage is used by
only binary solutions
 It is used for calculation of very
small quantities

15. Mole fraction
 It is defined as the ratio of number of
moles of a particular component to the
total number of moles of the solution. If
a substance ‘A’ dissolves in substance
‘B’ and their number of moles are nA
and nB respectively; then the mole
fractions of A and B are given as XA and
XB
 Mole fraction is a dimensionless
quantity

16. Molarity
 It is the most widely used unit and
is denoted by M. It is defined as
the number of moles of the solute
in 1 litre of the solution. Therefore
 molarity of a solution depends
upon temperature because
volume of a solution is

17. Molality
 It is defined as the
number of moles of solute
present in 1 kg of solvent.
It is denoted by m.
 molality of a solution does
not change with
temperature because

18. Two main concentrations
of solutions
molarity
 It is the most widely used
unit and is denoted by M. It
is defined as the number of
moles in solute upon
volume of solution in
litres(L).
 Molarity= no of moles in
solute/volume of solution in
litres.
 Molarity is not preferred
molality
 It is defined as the
number of moles of
solute present in 1 kg of
solvent. It is denoted by
(m).
 Molality= no of moles in
solute/ mass of solvent
in kg.
 Molality is preferred

19. Concepts in molarity
 In some cases to find the mass of a certain
compound volume of a compound , mass of the
second compound and volume of the second
compound will be given in that case we apply the
formula :-

20.  Another very important formula used in molarity is finding
molarity using density.
The formula is
M = % x d x10
mass
Example question:
Commercially available sulphuric acid contains 93% by mass and
has has a density of 1.84 g/cm3 . Calculate the molarity of the
solution and the volume of concentrated acid required to prepare
2.5mL of 0.50 M of H2SO4.

22. Examples on molality
What is the molality of a solution consisting of 1.34 mL of carbon tetrachloride
(CCl4, density= 1.59 g/mL) in 65.0 mL of methylene chloride (CH2Cl2, density
= 1.33 g/mL)?
Solution:-
1) Moles CCl4:
1.34 mL times 1.59 g/mL = 2.1306 g
2.1306 g / 153.823 g/mol = 0.013851 mol
2) Calculate the Mass of the methylene chloride:
65.0 mL times 1.33 g/mL = 86.45 g = 0.08645 kg
3) Calculate the Molality:
0.013851 mol / 0.08645 kg = 0.160 m